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Topic 5: Structure and Shape
Lewis structures
Lewis structures are a means of determining stable electron arrangements in molecules. It considers the valence electrons of an atom only. A stable arrangement is one in which each atom has achieved a Noble gas electron configuration by distribution of the electrons as bond pairs or lone pairs (non-bonded pairs). A Noble gas electron configuration is 2 for hydrogen and 8 for C, N, O and F. This is sometimes called The Octet Rule.
To draw a Lewis Structure you need to know which atoms are bonded to which.
Then:
- Add up the total number of valence electrons present, add or subtract electrons to account for any charge.
- Join the appropriate atoms using an electron pair for each bond.
- Distribute the remaining electrons to result in an octet of electrons on each atom (except hydrogen that always has two electrons associated with it).
- If there are too few electrons to give every atom an octet, move non-bonded pairs between atoms to give multiple bonds.
- If there are electrons left over after forming octets, place them on the central atom.
- Indicate the overall charge.
Question: Draw the Lewis structures of: CH 4 , NH 3 , NH 4 +, H 2 O, O 2 and N 2
Note: Be and B may have less than 8 electrons and Cl, Br, I, P, S may have more than 8 electrons in some compounds.
Question: Draw the Lewis structures of BeCl 2 and BF 3
Question: Draw the Lewis structures of PCl 3 , PCl 5 , SF 6 and ICl 4 -
Where more than one structure may be drawn, resonance occurs where the actual structure is a weighted combination of all possible structures and the electrons are delocalised in the molecule/ion.
Question: Draw the Lewis structure of O 3 , HCO 2 -^ , NO 3 -^ and SO 42 -
Molecular Geometry
The valence shell electron pair repulsion model (VSEPR model) assumes that electron pairs repel one another. This produces a set of geometries which depend only on the number of valence shell electron pairs and not on the atoms present.
To determine the molecular geometry
Draw the Lewis structure Count the number of electron pairs (bond pairs and lone pairs but count multiple bonds as one pair) Arrange electron pairs to minimise repulsion Position the atoms to minimise the lone pair - lone pair repulsion if > 1 lone pair Name the geometry from the atom positions
Question : Determine the molecular geometry of the following species:
CH 4
Geometry =
CH 3 -
Geometry =
NH 4 +
Geometry =
NH 3
Geometry =
NH 2 -
Geometry =
H 3 O+
Geometry =
H 2 O
Geometry =
Dipole Moments
Any bond between two different atoms will be polar as a result of the electronegativity difference between the atoms. A molecule has a permanent dipole moment if it contains polar bonds and is not a symmetrical shape.
Examples of polar molecules:
Examples of non-polar molecules:
Question: State whether the following molecules are polar or non-polar.
SO 3 SO 2 CH 4 SF 4 PCl 5 IF 5
N H (^) HH
O H Cl (^) H H
H C Cl (^) ClCl
C H H
O
O C O
Cl C Cl (^) ClCl
H H
F
S
F
F F
F F
Intra- & Inter-molecular forces
If matter is so attracted to itself, why doesn’t it squash down to nothing? There is a balance in energy between the forces of attraction and the strong force of repulsion that atoms experience when they get very close together.
The separation distance is known either as the: Bond length (for intramolecular cases), or the Van der Waals (VDW) distance (for intermolecular situatiuons - these are larger than bond lengths since the forces of attraction are weaker).
Intramolecular Forces
All attractive and repulsive interactions between atoms arise from electrical forces. We have already examined the bonding forces that hold atoms together. These are intramolecular bonding forces. One thing they have in common is that they are all quite STRONG (typically 80-4000 kJ mol -1^ ). When you break and form these “bonds” you are also performing a chemical reaction.
There are also intra- and inter-molecular non-bonding forces. These are largely WEAK (typically 0-50 kJ mol-1^ ) and when you make and break these you are mainly performing physical changes.
Charge distribution
Such interactions are all based on electrical forces. We recognise three different ways that the electron distribution can affect such interactions:
Ionic charge In this case the molecule itself has an excess of negative (anion) or positive (cation) charge.
Dipole moment Polar covalent bonds have an unequal sharing of electrons between the two atoms. In these cases one end of the bond is more negative and the other more positive. If the molecule is a diatomic species then we call the molecule “polar”. These “bond dipoles” can be added up in more complicated molecules. Frequently the sum of the bond dipoles gives rise to an overall dipole moment in a polyatomic molecule. However, sometimes, due to symmetry or accident, the bond dipoles can cancel to give rise to a non-polar molecule (even though the individual bonds are polar).
Polarizability The electrons on a molecule are never stationary, nor rigidly held. When a molecule is brought into the vicinity of other charges, the electrons on the molecule will move in response to this charge. The freedom of the electrons to move around is called the “polarizability” of the molecule. Small and first
Hydrogen bonds
Hydrogen bonding arises from an unusually strong and directed dipole-dipole force. When H is bonded to a very electronegative element (F, O, N) the bond is polar covalent. H is unusual because with only one electron, it leaves a partially exposed nucleus (H has no other core electrons to shield the nucleus). The bond can be thought of as forming between the hydrogen atom and the lone pairs of the F, N, or O in HF, NH 3 and H 2 O, respectively.
F, O and N are all 2nd^ row elements, which conveys certain properties: They are the most electronegative elements They are small (only 2 s , 2 p in outer shell) They have lone pairs
Water is perhaps the most unusual liquid. Each water molecule is H-bonded to FOUR other water molecules (donating 2 H-atoms and accepting two H-atoms to the lone pairs), forming a tetrahedral network (ice) and a loose tetrahedral network (liquid).
Measuring Intermolecular Forces
This can be done by measuring the molar enthalpy of vaporisation or other phase transition. An indirect measure is to measure the boiling point.
Question: Explain the trend in the table in terms of the type and size of IM Forces present.
Question: Which of the following substances show H-bonding? Draw the structure for those that do. a) C 2 H 6 b) CH 3 OH c) CH 3 CONH 2 d) H 2 S e) HCl f) NH 3
Substance ΔHvap (kJ mol -1^ ) H-bonds? M.W.
Butane, CH 3 (CH 2 ) 2 CH 3 22
Diethyl ether, (C 2 H 5 )-O-(C 2 H 5 ) 27
Methanol, CH 3 OH 38
Ethanol, CH 3 CH 2 OH 43
Water, H 2 O 44
Decane, CH 3 (CH 2 ) 8 CH 3 51
Blackman Figure 6.
Summary
Molecular Solids
The normal sequence of phases we expect to see on warming is Solid Liquid Gas
The melting transition from solid to liquid is not a good way of measuring intermolecular forces, as the molecules are held together by attractive forces in both phases. In order to get an idea of the strength of the interaction, we need to separate the molecules. This requires a transition to a gas, i.e. Liquid Gas (vaporisation or boiling) Solid Gas (sublimation)
Low boiling or sublimation points indicate weak intermolecular interactions, and high boiling points strong ones.
Melting points (Solid Liquid) can only be used as a rough guide, as the melting point must be below the boiling point.
1. Non-polar atoms and molecules Low boiling points E.g. Ar Tb = 87 K; N 2 Tb = 77 K; Propane Tb = 231 K 2. Polar molecules Broad range of boiling points depending on attractions. E.g. chloromethane (CH 3 Cl) Tb = 249 K; acetonitrile (CH 3 CN) Tb = 392 K. Boiling point increases with increasing dipole moment of molecule.
Silberberg Table 12.
Rigid rod-like molecules also form liquid crystals with positional as well as orientational order. These are called smectic phases, and form at lower temperatures than nematic phases.
In a smectic phase, the molecules are aligned but also organised in layers. This is due to packing of the rigid groups and also due to some entanglement and intercalation of the flexible alkyl chains.
In the smectic A phase the molecules are layered and oriented on average normal to the layers (shown on the right). In the smectic C phase the molecules are layered and more highly oriented at an angle to the layer normal.
The normal sequence (if all phases are present) is Solid Smectic C Smectic A Nematic Liquid
S C O O
CN
